| publisher=NASA | accessdate=2008-02-05 }}</ref> Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from [[hydrocarbon]]s such as methane, after which most elemental hydrogen is used "captively" (meaning locally at the production site), with the largest markets about equally divided between fossil fuel upgrading and ammonia production (mostly for the fertilizer market). Hydrogen may be produced from water using the process of electrolysis, but this process is presently significantly more expensive commercially than hydrogen production from natural gas.<ref>{{cite web

The most common naturally occurring [[isotope]] of hydrogen, known as [[hydrogen-1|protium]], has a single [[proton]] and no [[neutron]]s. In [[ionic compound]]s it can take on either a positive charge (becoming a [[Ion|cation]] composed of a bare proton) or a negative charge (becoming an [[Ion|anion]] known as a [[hydride]]). Hydrogen can form compounds with most elements and is present in [[water]] and most [[organic compound]]s. It plays a particularly important role in [[acid-base reaction theories|acid-base chemistry]], in which many reactions involve the exchange of protons between soluble molecules. As the only neutral atom for which the [[Schrödinger equation]] can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics.

+

+

==Chemistry and Characteristics==

+

The solubility and characteristics of hydrogen with various metals are very important in [[metallurgy]] (as many metals can suffer [[hydrogen embrittlement]]<ref>{{cite journal

+

| last=Rogers | first=H. C.

+

| title=Hydrogen Embrittlement of Metals

+

| journal=Science | year=1999

+

| volume=159 | issue=3819 | pages=1057-1064

+

| doi=10.1126/science.159.3819.1057 }}</ref>) and in developing safe ways to store it for use as a fuel.<ref>{{cite news

When mixed with oxygen across a wide range of proportions, hydrogen explodes upon ignition. Hydrogen burns violently in air. It ignites automatically at a temperature of 560&nbsp;°C.<ref>{{cite web

+

| url=http://physchem.ox.ac.uk/MSDS/HY/hydrogen.html

+

| title=Safety data for hydrogen | author=Staff

+

| accessdate=2008-02-05 | date=[[September 10]], [[2005]]

+

| work=Chemical and Other Safety Information

+

| publisher= The Physical and Theoretical Chemistry Laboratory, Oxford University }}</ref> Pure hydrogen-oxygen flames burn in the [[ultraviolet]] color range and are nearly invisible to the naked eye, as illustrated by the faintness of flame from the main [[Space Shuttle]] engines (as opposed to the easily visible flames from the shuttle boosters). Thus it is difficult to visually detect if a hydrogen leak is burning. The [[Hindenburg (airship)#Disaster|explosion of the Hindenburg airship]] was an infamous case of hydrogen combustion (pictured); the cause is debated, but combustible materials in the ship's skin were responsible for the coloring of the flames.<ref>{{cite web

+

| last = Dziadecki | first = John | year = 2005

+

| url = http://spot.colorado.edu/~dziadeck/zf/LZ129fire.htm

+

| title = Hindenburg Hydrogen Fire | accessdate = 2007-01-16 }}</ref> Another characteristic of hydrogen fires is that the flames tend to ascend rapidly with the gas in air, as illustrated by the Hindenburg flames, causing less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many of the deaths which occurred were from falling or from diesel fuel burns.<ref>{{cite web

+

| last=Werthmüller | first=Andreas

+

| url=http://www.hydropole.ch/Hydropole/Intro/Hindenburg.htm

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| title=The Hindenburg Disaster

+

| publisher=Swiss Hydrogen Association

+

| accessdate=2008-02-05 }}</ref>

+

+

H<sub>2</sub> reacts directly with other oxidizing elements. A violent and spontaneous reaction can occur at room temperature with [[chlorine]] and [[fluorine]], forming the corresponding hydrogen halides: [[hydrogen chloride]] and [[hydrogen fluoride]].<ref>{{cite book

+

| title=Handbook of Isotopes in the Cosmos: Hydrogen to Gallium

+

| last=Clayton | first=Donald D. | year=2003

+

| publisher=Cambridge University Press

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| isbn=0521823811 }}</ref>

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+

===Electron energy levels===

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{{main|Hydrogen atom}}

+

[[Image:hydrogen atom.svg|thumb|200px|right|Depiction of a hydrogen atom showing the diameter as about twice the [[Bohr model]] radius. (Image not to scale)]]

+

+

The [[ground state]] [[energy level]] of the electron in a hydrogen atom is -13.6 [[Electronvolt|eV]], which is equivalent to an ultraviolet [[photon]] of roughly 92 [[metre|nm]].<ref>{{cite web

+

| url=http://jupiter.phy.umist.ac.uk/~tjm/ISPhys/l7/ispl7.html

+

| title=Lecture 7, Emission Lines&nbsp;— Examples

+

| accessdate=2008-02-05 | last=Millar | first=Tom

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| date=[[December 10]], [[2003]]

+

| work=PH-3009 (P507/P706/M324) Interstellar Physics

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| publisher=University of Manchester }}</ref>

+

+

The energy levels of hydrogen can be calculated fairly [[accuracy|accurately]] using the [[Bohr model]] of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the sun. However, the [[electromagnetic]] force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other by [[gravity]]. Because of the discretization of [[angular momentum]] postulated in early [[quantum mechanics]] by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies.<ref>{{cite web

+

| last=Stern | first=David P. | date=[[May 16]], [[2005]]

+

| url=http://www-spof.gsfc.nasa.gov/stargaze/Q5.htm

+

| title=The Atomic Nucleus and Bohr's Early Model of the Atom

+

| publisher=NASA Goddard Space Flight Center

+

| accessdate=2007-12-20 }}</ref>

+

+

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the [[Schrödinger equation]] or the equivalent [[Feynman]] [[path integral formulation]] to calculate the [[probability amplitude|probability density]] of the electron around the proton.<ref>{{cite web

| accessdate=2008-02-05 }}</ref> In the [[orthohydrogen]] form, the spins of the two protons are parallel and form a triplet state; in the [[parahydrogen]] form the spins are antiparallel and form a singlet. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".<ref name="Tikhonov">{{cite journal

+

| last=Tikhonov | first=Vladimir I.

+

| coauthors=Volkov, Alexander A.

+

| title=Separation of Water into Its Ortho and Para Isomers

+

| journal=Science

+

| year=2002 | volume=296 | issue=5577 | pages=2363

+

| doi=10.1126/science.1069513

+

}}</ref> The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since the ortho form is an [[excited state]] and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The physical properties of pure parahydrogen differ slightly from those of the normal form.<ref name="NASA">{{cite web

| publisher=NASA | accessdate=2008-02-05 }}</ref> The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and [[methylene]].

+

+

The uncatalyzed interconversion between para and ortho H<sub>2</sub> increases with increasing temperature; thus rapidly condensed H<sub>2</sub> contains large quantities of the high-energy ortho form that convert to the para form very slowly.<ref>{{cite journal

}}</ref> The ortho/para ratio in condensed H<sub>2</sub> is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is [[exothermic]] and produces enough heat to evaporate the hydrogen liquid, leading to loss of the liquefied material. [[Catalyst]]s for the ortho-para interconversion, such as [[iron]] compounds, are used during hydrogen cooling.<ref name="Svadlenak">{{cite journal

A molecular form called [[protonated molecular hydrogen]], or H<sub>3</sub><sup>+</sup>, is found in the [[interstellar medium]] (ISM), where it is generated by ionization of molecular hydrogen from [[cosmic ray]]s. It has also been observed in the upper atmosphere of the planet [[Jupiter]]. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H<sub>3</sub><sup>+</sup> is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.<ref>{{cite web

+

| author=McCall Group, Oka Group | date=[[April 22]], [[2005]]

+

| url=http://h3plus.uiuc.edu/ | title=H3+ Resource Center

+

| publisher=Universities of Illinois and Chicago

+

| accessdate=2008-02-05 }}</ref>

+

+

===Compounds===

+

{{further|[[:Category:Hydrogen compounds|Hydrogen compounds]]}}

+

+

====Covalent and organic compounds====

+

While H<sub>2</sub> is not very reactive under standard conditions, it does form compounds with most elements. Millions of [[hydrocarbon]]s are known, but they are not formed by the direct reaction of elementary hydrogen and carbon (although [[synthesis gas]] production followed by the [[Fischer-Tropsch process]] to make hydrocarbons comes close to being an exception, as this begins with coal and the elemental hydrogen is generated in situ).{{Fact|date=March 2008}} Hydrogen can form compounds with elements that are more [[electronegative]], such as [[halogen]]s (e.g., F, Cl, Br, I); in these compounds hydrogen takes on a partial positive charge.<ref>{{cite web| last = Clark| first = Jim| title = The Acidity of the Hydrogen Halides| work = Chemguide| date = 2002| url = http://www.chemguide.co.uk/inorganic/group7/acidityhx.html#top| accessdate = 2008-03-09}}</ref> When bonded to [[fluorine]], [[oxygen]], or [[nitrogen]], hydrogen can participate in a form of strong noncovalent bonding called [[hydrogen bond]]ing, which is critical to the stability of many biological molecules.<ref>{{cite web| last = Kimball| first = John W. | title = Hydrogen| work = Kimball's Biology Pages| date = 2003-08-07| url = http://users.rcn.com/jkimball.ma.ultranet/BiologyPages/H/HydrogenBonds.html| accessdate = 2008-03-04}}</ref><ref>IUPAC Compendium of Chemical Terminology, Electronic version, [http://goldbook.iupac.org/H02899.html Hydrogen Bond]</ref> Hydrogen also forms compounds with less electronegative elements, such as the [[metal]]s and [[metalloid]]s, in which it takes on a partial negative charge. These compounds are often known as [[hydride]]s.<ref>{{cite web| last = Sandrock| first = Gary| title = Metal-Hydrogen Systems| publisher = Sandia National Laboratories| date = 2002-05-02| url = http://hydpark.ca.sandia.gov/DBFrame.html| accessdate = 2008-03-23}}</ref>

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+

Hydrogen forms a vast array of compounds with [[carbon]]. Because of their general association with living things, these compounds came to be called [[organic compound]]s;<ref name="hydrocarbon">{{cite web| title = Structure and Nomenclature of Hydrocarbons | publisher = Purdue University| url = http://chemed.chem.purdue.edu/genchem/topicreview/bp/1organic/organic.html| accessdate = 2008-03-23}}</ref> the study of their properties is known as [[organic chemistry]]<ref>{{cite web| title = Organic Chemistry| work = Dictionary.com| publisher = Lexico Publishing Group| date = 2008| url = http://dictionary.reference.com/browse/organic%20chemistry| accessdate = 2008-03-23}}</ref> and their study in the context of living [[organism]]s is known as [[biochemistry]].<ref>{{cite web| title = Biochemistry| work = Dictionary.com| publisher = Lexico Publishing Group| date = 2008| url = http://dictionary.reference.com/browse/biochemistry| accessdate = 2008-03-23}}</ref> By some definitions, "organic" compounds are only required to contain carbon (as a classic historical example, [[urea]]). However, most of them also contain hydrogen, and since it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry.<ref name="hydrocarbon"/>

+

+

In [[inorganic chemistry]], hydrides can also serve as [[bridging ligand]]s that link two metal centers in a [[coordination complex]]. This function is particularly common in [[group 13 element]]s, especially in [[borane]]s ([[boron]] hydrides) and [[aluminium]] complexes, as well as in clustered [[carborane]]s.<ref name="Miessler" />

+

+

====Hydrides====

+

Compounds of hydrogen are often called [[hydride]]s, a term that is used fairly loosely. To chemists, the term "hydride" usually implies that the H atom has acquired a negative or anionic character, denoted H<sup>−</sup>. The existence of the hydride anion, suggested by G.N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten [[lithium hydride]] (LiH), that produced a [[stoichiometric]] quantity of hydrogen at the anode.<ref name="Moers">{{cite journal

+

| last=Moers | first=Kurt

+

| title=Investigations on the Salt Character of Lithium Hydride

+

| journal=Zeitschrift für Anorganische und Allgemeine Chemie

+

| year=1920 | volume=113 | issue=191 | pages=179-228

+

}}</ref> For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH<sub>2</sub>, which is polymeric. In [[lithium aluminium hydride]], the AlH<sub>4</sub><sup>−</sup> anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride.<ref name="Downs">{{cite journal

Oxidation of H<sub>2</sub> formally gives the [[proton]], H<sup>+</sup>. This species is central to discussion of [[acid]]s, though the term proton is used loosely to refer to positively charged or [[cation]]ic hydrogen, denoted H<sup>+</sup>. A bare proton H<sup>+</sup> cannot exist in solution because of its strong tendency to attach itself to atoms or molecules with electrons. To avoid the convenient fiction of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain the [[hydronium]] ion (H<sub>3</sub>O<sup>+</sup>) organized into clusters to form H<sub>9</sub>O<sub>4</sub><sup>+</sup>.<ref name="Okumura">{{cite journal

}}</ref> Other [[oxonium]] ions are found when water is in solution with other solvents.<ref name="Perdoncin">{{cite journal

+

| last=Perdoncin | first=Giulio | coauthors=Scorrano, Gianfranco

+

| title=Protonation Equilibria in Water at Several Temperatures of Alcohols, Ethers, acetone, Dimethyl Sulfide, and Dimethyl Sulfoxide

+

| journal=Journal of the American Chemical Society

+

| year=1977 | volume=99 | issue=21 | pages=6983-6986

+

| doi=10.1021/ja00463a035

+

}}</ref>

+

+

Although exotic on earth, one of the most common ions in the universe is the [[Protonated molecular hydrogen|H<sub>3</sub><sup>+</sup>]] ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.<ref name="Carrington">{{cite journal

[[Image:Hydrogen.svg|thumb|150px|left|Protium, the most common isotope of hydrogen, has one proton and one electron. Unique among all stable isotopes, it has no neutrons. (see [[diproton]] for discussion of why others do not exist)]]

+

+

Hydrogen has three naturally occurring isotopes, denoted <sup>1</sup>H, <sup>2</sup>H, and <sup>3</sup>H. Other, highly unstable nuclei (<sup>4</sup>H to <sup>7</sup>H) have been synthesized in the laboratory but not observed in nature.<ref name="Gurov">{{cite journal

| title=Experimental Evidence for the Existence of <sup>7</sup>H and for a Specific Structure of <sup>8</sup>He

+

| journal=Physical Review Letters

+

| year=2003 | volume=90 | issue=8 | pages=082501

+

| doi=10.1103/PhysRevLett.90.082501 }}</ref>

+

* '''<sup>1</sup>H''' is the most common hydrogen isotope with an abundance of more than 99.98%. Because the [[atomic nucleus|nucleus]] of this isotope consists of only a single [[proton]], it is given the descriptive but rarely used formal name ''protium''.<ref>{{cite journal

+

| last=Urey | first=Harold C.

+

| coauthors=Brickwedde, F. G.; Murphy, G. M.

+

| title=Names for the Hydrogen Isotopes

+

| journal=Science | year=1933 | volume=78

+

| issue=2035 | pages=602-603

+

| url=http://www.sciencemag.org/cgi/content/citation/78/2035/602

+

| accessdate=2008-02-20 }}</ref>

+

* '''<sup>2</sup>H''', the other stable hydrogen isotope, is known as ''[[deuterium]]'' and contains one proton and one [[neutron]] in its nucleus. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called [[heavy water]]. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for <sup>1</sup>H-[[NMR spectroscopy]].<ref>{{cite journal

Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used). The symbols D and T (instead of <sup>2</sup>H and <sup>3</sup>H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for [[phosphorus]] and thus is not available for protium.<ref>{{cite web

+

| last=Krogt | first=Peter van der | date=[[May 5]], [[2005]]

+

| url=http://elements.vanderkrogt.net/elem/h.html

+

| publisher=Elementymology & Elements Multidict

+

| title=Hydrogen | accessdate=2008-02-20 }}</ref> In its [[IUPAC nomenclature|nomenclatural]] guidelines, the [[International Union of Pure and Applied Chemistry]] allows any of D, T, <sup>2</sup>H, and <sup>3</sup>H to be used, although <sup>2</sup>H and <sup>3</sup>H are preferred.<ref>§ IR-3.3.2, [http://www.iupac.org/reports/provisional/abstract04/connelly_310804.html Provisional Recommendations], Nomenclature of Inorganic Chemistry, Chemical Nomenclature and Structure Representation Division, IUPAC. Accessed on line [[October 3]], [[2007]].</ref>

Hydrogen is the most [[Natural abundance|abundant]] element in the universe, making up 75% of [[Baryon|normal matter]] by [[mass]] and over 90% by number of atoms.<ref>{{cite web

+

| first=Steve | last=Gagnon

+

| url=http://education.jlab.org/itselemental/ele001.html

+

| title=Hydrogen | publisher=Jefferson Lab

+

| accessdate=2008-02-05 }}</ref> This element is found in great abundance in [[star]]s and [[gas giant]] planets. [[Molecular cloud]]s of H<sub>2</sub> are associated with [[star formation]]. Hydrogen plays a vital role in powering [[stars]] through [[proton-proton reaction]] and [[CNO cycle]] [[nuclear fusion]].<ref>{{cite web

+

| last=Haubold | first=Hans | coauthors=Mathai, A. M.

+

| date=[[November 15]], [[2007]]

+

| url=http://www.columbia.edu/~ah297/unesa/sun/sun-chapter4.html

+

| title=Solar Thermonuclear Energy Generation

+

| publisher=Columbia University | accessdate=2008-02-12

+

}}</ref>

+

+

Throughout the universe, hydrogen is mostly found in the [[atomic]] and [[Plasma (physics)|plasma]] states whose properties are quite different from molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the [[sun]] and other [[stars]]). The charged particles are highly influenced by magnetic and electric fields. For example, in the [[solar wind]] they interact with the Earth's [[magnetosphere]] giving rise to [[Birkeland current]]s and the [[Aurora (phenomenon)|aurora]]. Hydrogen is found in the neutral atomic state in the [[Interstellar medium]]. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the [[Universe]] up to [[redshift]] ''z''=4.<ref>{{cite journal

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H<sub>2</sub> (for data see table). However, hydrogen gas is very rare in the Earth's atmosphere (1 [[part per million|ppm]] by volume) because of its light weight, which enables it to [[atmospheric escape|escape from Earth's gravity]] more easily than heavier gases. Still, hydrogen is the third most abundant element on the Earth's surface.<ref name="ArgonneBasic">{{cite web

+

| author=Dresselhaus, Mildred et al | date=[[May 15]], [[2003]]

+

| url=http://www.sc.doe.gov/bes/hydrogen.pdf | format=PDF

+

| title=Basic Research Needs for the Hydrogen Economy

+

| publisher=Argonne National Laboratory, U.S. Department of Energy, Office of Science Laboratory

+

| accessdate=2008-02-05

+

}}</ref> Most of the Earth's hydrogen is in the form of [[chemical compound]]s such as [[hydrocarbon]]s and [[water]].<ref name="Miessler">{{cite book

+

| first=Gary L. | last=Miessler | coauthors=Tarr, Donald A.

+

| year=2003 | title=Inorganic Chemistry | edition=3rd edition

+

| publisher=Prentice Hall | isbn=0130354716 }}</ref> Hydrogen gas is produced by some [[bacteria]] and [[algae]] and is a natural component of [[flatus]]. [[Methane]] is a hydrogen source of increasing importance.

H<sub>2</sub> is a product of some types of [[Fermentation (biochemistry)|anaerobic metabolism]] and is produced by several [[microorganism]]s, usually via reactions [[catalysis|catalyzed]] by [[iron]]- or [[nickel]]-containing [[enzyme]]s called [[hydrogenase]]s. These enzymes catalyze the reversible [[redox]] reaction between H<sub>2</sub> and its component two protons and two electrons. Creation of hydrogen gas occurs in the transfer of reducing equivalents produced during [[pyruvate]] [[fermentation (biochemistry)|fermentation]] to water.<ref>{{cite book

+

| first=Richard | last=Cammack

+

| coauthors=Robson, R. L. | year=2001

+

| title=Hydrogen as a Fuel: Learning from Nature

+

| publisher=Taylor & Francis Ltd

+

| isbn=0415242428 }}</ref>

+

+

[[Water splitting]], in which water is decomposed into its component protons, electrons, and oxygen, occurs in the [[light reaction]]s in all [[photosynthetic]] organisms. Some such organisms&nbsp;— including the [[alga]] ''[[Chlamydomonas reinhardtii]]'' and [[cyanobacteria]]&nbsp;— have evolved a second step in the [[dark reaction]]s in which protons and electrons are reduced to form H<sub>2</sub> gas by specialized hydrogenases in the [[chloroplast]].<ref>{{cite journal

| doi=10.1074/jbc.M503840200 }}</ref> Efforts have been undertaken to genetically modify cyanobacterial hydrogenases to efficiently synthesize H<sub>2</sub> gas even in the presence of oxygen.<ref>{{cite web

+

| first=H. O. | last=Smith | coauthors=Xu, Q | year=2005

+

| url=http://ec.europa.eu/food/fs/sfp/addit_flavor/flav15_en.pdf

+

| format=PDF

+

| title=IV.E.6 Hydrogen from Water in a Novel Recombinant Oxygen-Tolerant Cyanobacteria System

Hydrogen poses a number of hazards to human safety, from potential [[detonation]]s and [[fire]]s when mixed with [[air]] to being an [[Asphyxia|asphyxant]] in its pure, [[oxygen]]-free form.

+

<ref name=NASAH2>{{cite web

+

| authors=Brown, W. J. et al

+

| first=H. O. | last=Smith | coauthors=Xu, Q

+

| url=http://www.hq.nasa.gov/office/codeq/doctree/canceled/871916.pdf

+

| format=PDF | year=1997

+

| title=Safety Standard for Hydrogen and Hydrogen Systems

+

| publisher=[[NASA]] | accessdate=2008-02-05 }}</ref> In addition, [[liquid hydrogen]] is a [[cryogen]] and presents dangers (such as [[frostbite]]) associated with very cold liquids. Hydrogen dissolves in some metals, and, in addition to leaking out, may have adverse effects on them, such as [[hydrogen embrittlement]].<ref>{{cite web| last = Roberge| first = Pierre R.| title = Corrosion Doctors&nbsp;— Hydrogen Embrittlement| url = http://www.corrosion-doctors.org/Forms-HIC/embrittlement.htm| accessdate = 03-04-2008}}</ref> Hydrogen gas leaking into external air may spontaneously ignite. However, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to [[accident]]al [[burn]]s.<ref>{{cite web| title = Hydrogen Safety| publisher = Humboldt State University| url = http://www.humboldt.edu/~serc/h2safety.html| accessdate = 2008-03-15}}</ref>

+

+

Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the [[Spin isomers of hydrogen|parahydrogen/orthohydrogen]] ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen [[detonation]] parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.<ref name=NASAH2/>

*Logothetis, D. E., Boulos, Z., & Terman, M. (1984). Lick rate and the circadian rhythm of water intake in the rat: Effects of deuterium oxide: Annals of the New York Academy of Sciences Vol 423 1984, 614-617.

Hydrogen is the most abundant of the chemical elements, constituting roughly 75% of the universe's elemental mass.[2] Stars in the main sequence are mainly composed of hydrogen in its plasma state. Elemental hydrogen is relatively rare on Earth, and is industrially produced from hydrocarbons such as methane, after which most elemental hydrogen is used "captively" (meaning locally at the production site), with the largest markets about equally divided between fossil fuel upgrading and ammonia production (mostly for the fertilizer market). Hydrogen may be produced from water using the process of electrolysis, but this process is presently significantly more expensive commercially than hydrogen production from natural gas.[3]

The most common naturally occurring isotope of hydrogen, known as protium, has a single proton and no neutrons. In ionic compounds it can take on either a positive charge (becoming a cation composed of a bare proton) or a negative charge (becoming an anion known as a hydride). Hydrogen can form compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry, in which many reactions involve the exchange of protons between soluble molecules. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics.

Hydrogen gas is highly flammable and will burn at concentrations as low as 4% H2 in air.[9] The enthalpy of combustion for hydrogen is −286 kJ/mol;[10] it burns according to the following balanced equation.

When mixed with oxygen across a wide range of proportions, hydrogen explodes upon ignition. Hydrogen burns violently in air. It ignites automatically at a temperature of 560 °C.[12] Pure hydrogen-oxygen flames burn in the ultraviolet color range and are nearly invisible to the naked eye, as illustrated by the faintness of flame from the main Space Shuttle engines (as opposed to the easily visible flames from the shuttle boosters). Thus it is difficult to visually detect if a hydrogen leak is burning. The explosion of the Hindenburg airship was an infamous case of hydrogen combustion (pictured); the cause is debated, but combustible materials in the ship's skin were responsible for the coloring of the flames.[13] Another characteristic of hydrogen fires is that the flames tend to ascend rapidly with the gas in air, as illustrated by the Hindenburg flames, causing less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many of the deaths which occurred were from falling or from diesel fuel burns.[14]

Electron energy levels

The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the sun. However, the electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies.[17]

Elemental molecular forms

There are two different types of diatomic hydrogen molecules that differ by the relative spin of their nuclei.[19] In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state; in the parahydrogen form the spins are antiparallel and form a singlet. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the "normal form".[20] The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but since the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The physical properties of pure parahydrogen differ slightly from those of the normal form.[21] The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene.

The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that convert to the para form very slowly.[22] The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate the hydrogen liquid, leading to loss of the liquefied material. Catalysts for the ortho-para interconversion, such as iron compounds, are used during hydrogen cooling.[23]

A molecular form called protonated molecular hydrogen, or H3+, is found in the interstellar medium (ISM), where it is generated by ionization of molecular hydrogen from cosmic rays. It has also been observed in the upper atmosphere of the planet Jupiter. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H3+ is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium.[24]

Compounds

Covalent and organic compounds

While H2 is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon (although synthesis gas production followed by the Fischer-Tropsch process to make hydrocarbons comes close to being an exception, as this begins with coal and the elemental hydrogen is generated in situ).[How to reference and link to summary or text] Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I); in these compounds hydrogen takes on a partial positive charge.[25] When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules.[26][27] Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.[28]

Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds;[29] the study of their properties is known as organic chemistry[30] and their study in the context of living organisms is known as biochemistry.[31] By some definitions, "organic" compounds are only required to contain carbon (as a classic historical example, urea). However, most of them also contain hydrogen, and since it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word "organic" in chemistry.[29]

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. To chemists, the term "hydride" usually implies that the H atom has acquired a negative or anionic character, denoted H−. The existence of the hydride anion, suggested by G.N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), that produced a stoichiometric quantity of hydrogen at the anode.[33] For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH2, which is polymeric. In lithium aluminium hydride, the AlH4− anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride.[34] Binary indium hydride has not yet been identified, although larger complexes exist.[35]

"Protons" and acids

Oxidation of H2 formally gives the proton, H+. This species is central to discussion of acids, though the term proton is used loosely to refer to positively charged or cationic hydrogen, denoted H+. A bare proton H+ cannot exist in solution because of its strong tendency to attach itself to atoms or molecules with electrons. To avoid the convenient fiction of the naked "solvated proton" in solution, acidic aqueous solutions are sometimes considered to contain the hydronium ion (H3O+) organized into clusters to form H9O4+.[36] Other oxonium ions are found when water is in solution with other solvents.[37]

Although exotic on earth, one of the most common ions in the universe is the H3+ ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.[38]

Isotopes

Hydrogen has three naturally occurring isotopes, denoted 1H, 2H, and 3H. Other, highly unstable nuclei (4H to 7H) have been synthesized in the laboratory but not observed in nature.[39][40]

1H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.[41]

2H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for 1H-NMR spectroscopy.[42] Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.[43]

Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were given names, but such names are no longer used). The symbols D and T (instead of 2H and 3H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium.[49] In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry allows any of D, T, 2H, and 3H to be used, although 2H and 3H are preferred.[50]

Throughout the universe, hydrogen is mostly found in the atomic and plasma states whose properties are quite different from molecular hydrogen. As a plasma, hydrogen's electron and proton are not bound together, resulting in very high electrical conductivity and high emissivity (producing the light from the sun and other stars). The charged particles are highly influenced by magnetic and electric fields. For example, in the solar wind they interact with the Earth's magnetosphere giving rise to Birkeland currents and the aurora. Hydrogen is found in the neutral atomic state in the Interstellar medium. The large amount of neutral hydrogen found in the damped Lyman-alpha systems is thought to dominate the cosmological baryonic density of the Universe up to redshiftz=4.[53]

Under ordinary conditions on Earth, elemental hydrogen exists as the diatomic gas, H2 (for data see table). However, hydrogen gas is very rare in the Earth's atmosphere (1 ppm by volume) because of its light weight, which enables it to escape from Earth's gravity more easily than heavier gases. Still, hydrogen is the third most abundant element on the Earth's surface.[54] Most of the Earth's hydrogen is in the form of chemical compounds such as hydrocarbons and water.[32] Hydrogen gas is produced by some bacteria and algae and is a natural component of flatus. Methane is a hydrogen source of increasing importance.
[55]

Safety and precautions

Hydrogen poses a number of hazards to human safety, from potential detonations and fires when mixed with air to being an asphyxant in its pure, oxygen-free form.
[60] In addition, liquid hydrogen is a cryogen and presents dangers (such as frostbite) associated with very cold liquids. Hydrogen dissolves in some metals, and, in addition to leaking out, may have adverse effects on them, such as hydrogen embrittlement.[61] Hydrogen gas leaking into external air may spontaneously ignite. However, hydrogen fire, while being extremely hot, is almost invisible, and thus can lead to accidentalburns.[62]

Even interpreting the hydrogen data (including safety data) is confounded by a number of phenomena. Many physical and chemical properties of hydrogen depend on the parahydrogen/orthohydrogen ratio (it often takes days or weeks at a given temperature to reach the equilibrium ratio, for which the data is usually given). Hydrogen detonation parameters, such as critical detonation pressure and temperature, strongly depend on the container geometry.[60]

Logothetis, D. E., Boulos, Z., & Terman, M. (1984). Lick rate and the circadian rhythm of water intake in the rat: Effects of deuterium oxide: Annals of the New York Academy of Sciences Vol 423 1984, 614-617.